Factors affecting rate of reaction

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Intros
Lessons
  1. Controlling rate of reaction
  2. What affects the rate?
  3. What is a catalyst?
  4. Why are catalysts and solutions important?
  5. How phase affects reaction rate.
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Examples
Lessons
  1. Apply knowledge of factors affecting reaction rate to chemical reactions.
    The reaction between sodium hydroxide, NaOH, and hydrogen chloride, HCl, is shown below:

    2NaOH+2HCL→2NaCltoH2O\mathrm{2NaOH + 2 HCL \to 2 NaCl to H_2O}

    1. Hydrogen chloride has a boiling point of -85°C and sodium hydroxide has a melting point of 318°C. Give a reason why this reaction is very slow with the reactants in their neutral state at room temperature.
    2. What change could a chemist make to increase the rate of this reaction? Explain why this change increases the rate of reaction.
  2. Apply knowledge of factors affecting reaction rate to chemical reactions.
    The decomposition of aqueous hydrogen peroxide is shown in the equation below. This reaction occurs at room temperature by itself but is slow:

    2H2O2(aq)→2H2O(l)+O2(g) \mathrm{ 2 H_2 O_{2(aq)} \to 2H_2O_{(l)} + O_{2(g)} }

    1. The rate of this reaction is increased dramatically by adding potassium iodide, KI, to the reaction vessel. Potassium iodide is not used in the reaction. What is the role of potassium iodide?
    2. Explain, in terms of activation energy, why adding potassium iodide speeds the reaction up.

      i) Explain, in terms of activation energy, why adding potassium iodide speeds the reaction up.

      ii) What else in the equation suggests a generally quicker rate of reaction?
  3. Apply knowledge of factors affecting reaction rate to chemical reactions.
    From the reactions listed below, identify which are homogeneous and compare their relative rates of reaction.

    i) Na(s)+Cl(g)→NaCl(s) \mathrm{Na_{(s)} + Cl_{(g)} \to NaCl_{(s)} }

    ii) CH3COOH(aq)+H2O(l)→CH3COO(aq)−+H3O(aq)+ \mathrm{ CH_3COOH_{(aq)} + H_2O_{(l)} \to CH_3COO^-_{(aq)} + H_3O^+_{(aq)}}

    iii) Mg(s)+H2O(l)→MgOH(aq)+H2(g) \mathrm{ Mg_{(s)} + H_2O_{(l)} \to MgOH_{(aq)} + H_{2(g)}}
    Topic Notes
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    In this lesson, we will learn:

    • To know and describe the effects of basic factors that affect the rate of reaction.
    • To explain the effect of a catalyst on the rate of a chemical reaction.
    • To apply ideas of surface area to explain the importance of catalysts and solutions to reaction rate.
    Notes:

    • Once scientists started measuring the rate of reactions; new areas of study developed that measured the factors that affected the rate of reaction. These are normally to do with the conditions the reaction is happening under.
    • There are many factors that affect the rate of reaction, which have been determined by experiment. Some are:
      • 1: Temperature: The higher the temperature of a reaction, the faster the reaction happens because the time taken for the reaction to happen decreases.
      • 2: Concentration: As concentration of reactants increases, time taken for the reaction decreases therefore the rate increases.
      • 3: Pressure: Pressure is like concentration for gaseous reactants – greater pressure forces gas particles together like in a high concentration solution. Therefore the higher the pressure, the quicker the rate of reaction.
      • 4: Surface area and catalysts: The greater surface area where particles can collide, the more particles will collide. Any way you can increase surface area for reactants to collide will decrease the time taken for the reaction to occur, and increase the reaction rate – the phase the reactants are in and catalysts both affect this.
    • The above four factors can affect any chemical reaction, but the rate will also depend on what chemical bonds need breaking (the reactant properties) in the reaction! Stronger bonds will require more energy to be overcome and so the reaction rates are naturally lower than chemical reactions where weak bonds are broken, or where very stable products are formed.
    • Catalysts affect the rate of reaction by providing an alternative reaction pathway of lower activation energy than the original uncatalyzed route. The catalyst itself remains unchanged in the reaction.
      • Think of a catalyst's effect like trying to climb a high wall without any help. Only the strongest and best climbers can do it!
      • If there was a ledge to climb onto first (the catalyst), climbing over the wall is a lot easier (the activation energy becomes lower) than before.
      • The ledge itself doesn't climb over and it doesn't change from being in the process – it just gets a little worn from it.
    • The reason that catalysts increase rate of reaction is that catalysts provide more surface area for reactant particles to collide and form the products. This is the same for why phase is important to reaction rate. Below is a summary of the importance of how phases affect reaction rate:
      • Chemical substances in the solid state have particles that are tightly packed and unable to move freely, so reactions between/involving solids are very slow.
      • In the liquid state, particles have more energy, are able to move and are close together so the reaction rate is increased.
      • In the gas state particles are highly energetic and a higher proportion of the collisions between reactant particles are successful collisions. Reaction rates between or involving gases are much higher than solids.
      • We saw in solution chemistry (C11.8.1) that solutions are important to use in chemical reactions and they behave as if they are their own state. This is because solutions:
        • Allow free movement of reactant particles throughout a liquid medium.
        • Allow reactants close proximity to each other.
        • Allows reactants to form aqueous ions.
        • Enables the positive/negative attractive forces to occur between reactant particles.
      • For these reasons reactions taking place in solutions are generally faster than in any other phase.
    • To summarize above: the relationship between rate of reaction and phase of reactants is as follows, fastest to slowest:

      Solution (aqueous) > Gases and liquids > Solids.

      This is observed by experiment and explainable by particle theory – the states where particles will collide with sufficient energy most frequently are the states with the fastest reaction rates.
    • There are two definitions given for reactions depending on which states are being used:
      • Heterogeneous reactions are reactions where the reactants are in different phases.
      • Homogeneous reactions are reactions where the reactants are all in the same phase. This includes all reactants dissolved in a solvent (even if the reactants were different phases before being dissolved!) and two liquids which completely dissolve in the other.
      • Because it is easier to control and manipulate the reaction vessel, homogeneous reactions are generally an advantage but aren't always an option.